Bio Midterm PDF

Midterm Chapter Two (Charles Despres) Matter - Matter is anything that has mass and occupies space - All life forms are composed of matter - Matter can exist in any of three states: solid, liquid, or ga Atoms - The smallest functional units of matter that form all chemical substances...

Midterm Chapter Two (Charles Despres) Matter - Matter is anything that has mass and occupies space - All life forms are composed of matter - Matter can exist in any of three states: solid, liquid, or ga Atoms - The smallest functional units of matter that form all chemical substances - Cannot be further broken down into other substances by ordinary means - Chemists study atoms and molecules, which are two or more atoms bonded together - Each specific type of atom is a chemical element Three subatomic particles Protons - positive charge (+) - found in atomic nucleus Neutrons - neutral - found in atomic nucleus Electrons - negative charge (−) - found in orbitals ❖ Protons and electrons are present in equal numbers, giving the atom no net charge ❖ The number of neutrons canvary Electrons occupy orbitals Scientists initially visualized an atom as a miniature solar system - This is an oversimplified but convenient image Electrons travel within regions surrounding the nucleus (orbitals) in which the probability of finding that electron is high. Better model of an atom is a central nucleus surrounded bycloudlike orbitals. Orbitals - S orbitals are spherical - P orbitals are propeller or dumbbell shaped - Each orbital can hold only 2 electrons Electron Shells Atoms with more electrons have orbitals within electron shells that are at greater and greater distances from the center of the nucleus 1st shell - 1 spherical orbital (1s) - holds one pair of electrons 2nd shell - 1 spherical orbital (2s) - holds one pair of electrons - 3 dumbbell-shaped orbitals (2p) - three pairs of electrons - Can hold four pairs of electrons = 8 electrons Example: Nitrogen atom 7 protons and 7 electrons 2 electrons fill 1st shell - 2 in the 1s orbital 5 electrons in 2nd shell - 2 fill the 2s orbital - 1 in each of the three 2p orbitals - Note: the outer 2nd shell is not full Electrons in the outer shell available to combine with other atoms are called valence electrons. a) Simplified depiction of a nitrogen atom Lewis Structures Here's an example ofNH3​(ammonia),​ where​you​can ​observe​the​lone​pair​ and​the​bonded ​pair. Lewis Structures Bonded pairs are shown as lines as it keeps things neat. Here's an example of H2O Protons Number of protons is what distinguishes one element from another Atomic number - Equals number of protons - Also equal to the number of electrons in the atom so that the net charge is zero Periodic table - Organized by atomic number. - Rows correspond to number of electron shells. - Columns, from left to right, indicate the numbers of electrons in the outer shell (the number of valence electrons) - Similar properties of elements within a column occur because they have the same number of electrons in their outer shells, and therefore they have similar chemical bonding properties. Atomic mass Protons and neutrons are nearly equal in mass, and both are more than 1,800 times the mass of an electron. Atomic mass scale indicates an atom’s mass relative to the mass of other atoms. Most common form of carbon has six protons and six neutrons, is assigned an atomic mass of exactly 12. - Hydrogen atom (atomic mass of 1) has 1/12 the mass of a carbon atom. - Magnesium atom (atomic mass of 24) has twice the mass of a carbon atom. Atomic mass in relation to mass of an electron Mass versus weight Weight is derived from the gravitational pull on a given mass A man weighs 154 pounds on Earth - On the moon he weighs about 25 pounds - On a neutron star’s surface he would weigh 21 trillion pounds His mass is the same in all locations Units Dalton - Unit of measurement for atomic mass - Also known as atomic mass unit (amu) - One Dalton (Da) equals 1/12 the mass of a carbon atom. - Carbon has an atomic mass of 12 Daltons Mole - 1 mole of any element contains the same number of atoms: 6.022x10^23 - Avogardo’s number Isotopes - Multiple forms of an element that differ in the number of neutrons - 12^C contains 6 protons and 6 neutrons - 14^C contains 6 protons and 8 neutrons - Atomic masses are averages of the masses of different isotopes of an element - Radioisotopes are unstable, emit radiations as the decay - Radioisotopes used in medicine for cancer treatment, imaging (PET scan) Hydrogen, oxygen, carbon, nitrogen Make up about 95% of the atoms in living organisms - Hydrogen and oxygen occur primarily in water - Nitrogen is found in proteins - Carbon is the building block of all living matter Mineral elements - less that 1% Trace elements - less that 0.01% - Yet they are essential for normal growth and function Chemical Bonds and Molecules Molecule - Two or more atoms bonded together Molecular formula - Contains chemical symbols of the elements in the molecule ( C6H12O6 ) - Subscript indicates how many of each atom are present (H20 has two hydrogens, 1 oxygen) Compound - Any molecule composed of two or more elements - H2O; C6H12O6 - Properties of a compound can be drastically different that the properties of the individual elements in the compound Three types of bonds Covalent bonds - Electrons are shared to fill valence shells - Can be polar covalent or nonpolar covalent Hydrogen Bond - Hydrogen atom from one polar molecule is attracted to an electronegative atom from another molecule Ionic Bond - Electrons are transferred, forming ions that are attracted to each other. Covalent Bonds - Atoms share a pair of electrons. - Occurs between atoms with unfilled outer electron shells - Covalent Bonds are strong chemical bonds because the shared electrons behave as if they belong to each atom. Can share… 1 pair of electrons - single bond 2 pairs of electrons - double bond 3 pairs of electrons - triple bond Octet Rule - Atoms are stable when their outer shell is full - For many atoms, the outer shell is filled with 8 electrons - An exception is hydrogen, which only has two electrons in it’s outer shell Polar covalent bonds - Form between atoms of different electronegativity (attraction to electrons) - Shared electrons are more likely to be close to the more electronegative atom - The unequal distribution of electrons create polarity(difference in electric charge ) across the molecule - Examples include O-H and N-H Nonpolar Covalent Bonds - Between atoms with similar electronegativities (attration to electrons) - Equal sharing of electrons - No charge difference across molecule - Examples include C-C and C-H Hydrogen Bonds The hydrogen atom from one polar molecule is attracted to an electronegative atom of another. Represented as dashed or dotted lines Individually, these are weak bonds that can form and break easily Collectively many H bonds can be strong overall - Holds DNA strands together Enzymes are molecules that catalyze biologically important chemical reactions - Small molecules may bind to enzymes via hydrogen bonds Van de Waals dispersion forces - Another type of weak molecular attraction - Arise because electrons are located within orbitals in a random way - A fleeting electrical attraction to other nearby molecule may arise - Collective strength can be quite strong Iconic Bonds An ion is an atom or molecule that has gained or lost one or more electrons Cations - have a net positive charger (+) Anions - have a net negative charge (-) Ionic bond occurs when a cation binds to an anion by electrostatic attraction Example is NaCl, table salt Molecules may change their shapes - Atoms combine to form a molecule with three dimensional shape - The shape is determined by the arrangement and number of bonds between atoms - Angles that form between atoms give molecules specific shapes - Covalent bonds are not rigid and rotation around single covalent bonds allows molecules to change shape. - The binding of one molecule to another can cause the molecule to change shape Free radicals - Molecule containing an atom with a single, unpaired electron in its outer shell - Highly reactive molecules; can “steal” an electron from other molecules - Can form by exposure to radiation and some toxins Examples are: O2, OH, NO - Can cause cell damage - Can kill invading bacteria - Benefits of antioxidants – protective compounds that can donate electrons without becoming highly reactive themselves. Chemical Reactions When one or more substances are changed into other substances Reactants -> products Properties of chemical reactions: - Require a source of energy - In living organisms, they often require an enzyme as catalyst but will eventually reach equilibrium. - Occur in liquid (water) 1.1 Properties of water - Bodies of all organisms largely composed of water - Up to 95% of the weight of certain plants comes from water - 60 to 70% of human body weight comes from water - Water is an important liquid in the surrounding environments of many organisms - Most chemical reactions in nature involve molecules that are dissolved in water, including reactions inside cells. Intracellular fluid - fluids inside the nucleus Extracellular fluid - fluids outside the nucleus 1.2 Properties of water Solution = solutes in a solvent - Solutes are dissolved substances - Solvent is the liquid - In an aqueous solution, water is the solvent - Ions and molecules with polar covalent bonds will dissolve in water - These are hydrophilic Solutes Hydrophilic - “water-loving” - Readily dissolve in water - Molecules with ionic and/or polar covalent bonds Hydrophobic - “water-fearing” - Do not dissolve in water - Nonpolar molecules like hydrocarbons, oils Amphipathic - “both loves” - Have both polar/ionized and nonpolar regions - May form micelles in water - Detergent is an amphipathic molecule - Polar (hydrophilic) regions at the surface of the micelle - Nonpolar ( hydrophobic) ends are oriented toward the interior of the micelle Measuring solutions Concentration - Amount of a solute dissolved in a unit volume of solution - 1 gram of NaCl dissolved in 1 liter of water = 1 gram/Liter Molarity - Molecular mass is the sum of all the atomic masses of all the atoms in the molecule - Molarity = Number of moles of a solute dissolved in 1 Liter of water - 1 mole of a substance is the amount of the substance in grams equal to its atomic or molecular mass H20 in three states of matter Solid (ice), liquid (water), and gas (water vapor) Changes in state, such as changes between the solid, liquid, and gas states of H2O, involve an input or release of energy - Heat of vaporization– energy to boil - Heat of fusion– energy to melt Specific heat is the amount of heat energy to raise temperature 1° Celsius Water is extremely stable as a liquid, due to high heats of vaporization and fusion, and high specific heat Colligative properties of water - Properties that depend strictly on the total number of dissolved solute particles, not on the type of solute. - Temperature at which a solution freezes or boils is influenced by amounts of dissolved solutes Addition of solutes to water - lowers the freezing point below 0° Celsius - raises the boiling point above 100° Celsius Antifreeze - (ethylene glycol) lowers the freezing point of the water and prevents it from freezing in cold weather Not just a solvent Water has many important functions in living organisms: - Participates in chemical reactions (hydrolysis or condensation) - Provides force or support - Removes toxic waste components - Evaporative cooling - Cohesion (molecules of the same type attract each other) and adhesion (unlike molecules attract each other) - Surface tension – measure of attraction between molecules at the surface of a liquid - Lubrication Acids and Bases1 - Pure water ionizes to a very small extent into hydrogen ions (H ) and hydroxide ions (OH ) Acids and Bases2 - Acids are molecules that release hydrogen ions in solution - A strong acid releases more H than a weak acid Bases lower the H+ concentration - Some release OH- - Others bind H+ The pH scale - pH= -𝑙𝑜𝑔10 (Η+) - Acidic solutions are pH 6 or below - pH 7 is neutral - Alkaline solutions are pH 8 or above Effects of pH The pH of a solution can affect - The shapes and functions of molecules - The rates of many chemical reactions - The ability of two molecules to bind to each other - The ability of ions or molecules to dissolve in water Buffers - Organisms usually tolerate only small changes in pH - Buffers help to maintain a constant pH - An acid-base buffer system can shift to remove or release H+ to adjust for changes in pH BIOLOGY CHAPTER 3 The Carbon Atom - Organic molecules contain carbon - Organic molecules are abundant in living organisms - Macromolecules are large, complex organic molecules Organic Chemistry  - Science of carbon-containing molecules - Vitalism - 19th century concept that organic molecules were created by and imparted with a vital life force within a plant or animal’s body - Believed organic compounds could not be synthesized - Later disproven – organic compounds can be synthesized Carbon Carbon has 4 electrons in its outer shell Needs 4 more electrons to fill the shell It can make up to four bonds - Usually single or double bonds Carbon can form nonpolar or polar bonds - Molecules with polar bonds are water soluble - Molecules with nonpolar bonds (like hydrocarbons) are not very water-soluble Functional Groups - Groups of atoms with special chemical features that are functionally important - Each type of functional group exhibits the same properties in all molecules in which it occurs Isomers Two molecules with an identical molecular formula but different structures and characteristics Structural isomers - contain the same atoms but in different bonding relationships Stereoisomers - identical bonding relationships, but the spatial positioning of the atoms differs in the two isomers - Cis-trans isomers: positioning around double bond - Enantiomers: mirror image molecules Difference in orientation leads to different binding abilities Enzymes that recognize one enantiomer usually do not recognize the other Synthesis and Breakdown of Organic Molecules and Macromolecules - Some organic molecules are large macromolecules composed of thousands or millions of atoms - Formed by linking monomers (one part) and polymers (many parts) - When a polymer is formed, two smaller molecules combine through a condensation reaction – produces a larger organic molecule plus a water molecule Polymer formation by dehydration (condensation) reactions - A molecule of water is removed each time a new monomer is added, thus a “dehydration” reaction - The process repeats to form long polymers - A polymer can consist of thousands of monomers - Dehydration is catalyzed by enzymes Breakdown of a polymer by hydrolysis reactions - A molecule of water is added back each time a monomer is released - The process repeats to break down long polymer - Hydrolysis is catalyzed by enzymes Four Major Classes of Organic Molecules Found in Living Cells - Carbohydrates - Lipids - Proteins - Nucleic acids Carbohydrates - Composed of carbon, hydrogen, and oxygen atoms - Cn(H2O)n - Most of the carbon atoms in a carbohydrate are linked to a hydrogen atom and a hydroxyl group Monosaccharides ❖ Simplest sugars Most common are 5 or 6 carbons - Pentoses - Ribose C5H10O5 - Deoxyribose (C5H10O4) Hexose  - Glucose (C6H12O6) Different ways to depict structures - Ring - Linear Glucose isomers Stereoisomers of glucose a- and β glucose - Hydroxyl group of carbon 1 is above or below ring D- and L-glucose - Enantiomers with mirror image structure - D-glucose commonly found in living cells - L-glucose rarely found in living cells Galactose - Hydroxyl group on carbon 4 of glucose is above the plane of the ring instead of below it Disaccharides - Composed of two monosaccharides - Joined by dehydration or condensation reaction ❖ Glycosidic bond Examples: sucrose, maltose, lactose Polysaccharides - Many monosaccharides linked together to form long polymers Examples: - Energy storage – starch, glycogen - Structural – cellulose, chitin, glycosaminoglycans, peptidoglycan Lipids - Composed predominantly of hydrogen and carbon atoms, and some oxygen - The defining feature of lipids is that they are nonpolar and therefore very insoluble in water  - Include fats, phospholipids, steroids, waxes - Lipids comprise about 40% of the organic matter in the average human body Fats1 - Also known as triglycerides - Formed by bonding glycerol to 3 fatty acids - Joined by dehydration; the resulting bond is an ester bond Chapter 6 Midterm: October 21st 1. Energy & Chemical reactions Chemical reactions: the process in which one or more substances are changed into other substances. - Chemical reactions occur when atoms combine with or dissociate from other atoms - Chemical bonds are energy relationships involving interactions among the electrons of reacting atoms. 4H + C -> CH4 (methane) - Metabolism is the total of all chemical reactions that occur within an organism Matter and Energy Matter: anything that occupies space and has mass and has 3 states - Solid: definite shape and volume - Liquid: definite volume, shape of container Energy: the ability to do work - Has no mass and does not take up space - Kinetic energy: energy is doing work - Potential energy: energy is inactive or stored Forms of energy: - Chemical energy is stored in chemical bonds of substances - Electrical energy results from - Mechanical energy is energy directly involved in moving matter - Radiant energy travels in waves, energy of the electromagnetic spectrum Two forms: - Kinetic energy: associated with movement - Potential energy: Energy held by an object because of its position relative to other objects Thermodynamics Thermodynamics: study of energy interconversions ( energy being converted from one form to another) ❖ First Law of Thermodynamics: - Law of conservation of energy - Energy cannot be created or destroyed but can be transformed from one type to another (chemical energy = heat ) ❖ Second Law of thermodynamics: - Transfer of energy from one form to another increases the entropy (degree of disorder) of a system - As entropy increases, less energy is available for organisms to use to promote change or do work (unusable energy) Change in free energy determines Direction of chemical reactions - Total energy= Usable energy + Unusable energy - Energy transformations involve an increase in entropy (a disorder that cannot be harnessed to do work) - Free energy(G)= amount of energy available to do work Change in free energy determines direction of chemical reactions H= G + TS - H = enthalpy or total energy - G = free energy or amount of energy for work - S = entropy or unusable energy - T = absolute temperature in kelvin (K) Spontaneous Reactions: - Occur without input of additional energy (will) proceed naturally - You may have to provide some activation energy, the rest will proceed without the need for continuous input of an external energy source. - E.g. Combustion Non-spontaneous reactions: A continuous energy input is necessary for the reaction to proceed - E.g. Photosynthesis is a process in that plants use to make glucose - Photosynthesis requires energy  - At night when it’s dark, photosynthesis does not work How to tell? deltaG = deltaH-TdeltaS - Key factor is the free energy change - if delta G is negative, then the process in spontaneous - Exergonic = spontaneous - delta G < 0 (negative free energy change) ( energy is released by reaction) - Endergonic = not spontaneous - Delta G > 0 (positive free energy change) ( requires the addition of energy to drive reaction) Hydrolysis of ATP: Spontaneous or non-spontaneous ΔG= −7.3kcal/mole - Exergonic = spontaneous: ΔG < 0 (negative free energy change) - Endergonic = not spontaneous: ΔG > 0 (positive free energy change) - Reaction favors formation of products - The energy liberated is used to drive a variety of cellular processes Cells use ATP hydrolysis to drive reactions - An endergonic reaction can be coupled to an exergonic reaction so that the two reaction overall is thermodynamically favored! - ATP is the major 'energy' molecule produced by metabolism: it’s “dispatched” to wherever a non-spontaneous reaction needs to occurs within the cell - The reactions will be spontaneous if the net free energy change for both processes is negative ATP Drives endergonic reactions ΔG = + 3.3Kcal/mole (endergonic) + ΔG = − 7.3Kcal/mole (exergonic) =  ΔG = − 4.0Kcal/mole (exergonic) Coupled reaction = spontaneous! Cells use ATP hydrolysis to drive reactions - Endergonic reactions can be coupled with exergonic reactions so the reaction overall is thermodynamically favored - The reactions will be spontaneous if the net free energy change for both processes is negative - In the couple reaction a phosphate is directly transferred from ATP to glucose - phosphorylation - Typical cell uses millions of ATP molecules per second to drive endergonic processes - Breakdown of food releases energy that allows cells to make more ATP to ADP Enzymes - Each ATP undergoes 10,000 cycles of hydrolysis and resynthesis every day - Particular amino acid sequences in proteins function as ATP binding sites - On average, 20% of all protein bind ATP - Likely an underestimate because there may be other types of ATP- binding sites - This illustrates the enormous importance of ATP as an energy source  - A spontaneous reaction is not necessarily a fast reaction - Catalyst an agent that speeds up the rate of a chemical reaction without being consumed during the reaction - Enzymes protein catalysts in living cells - Ribozymes: RNA molecules with catalytic properties - Enzymes - Act as biological catalysts - Increase the rate of chemical reactions - Bind to substrates at an active site to catalyze reactions - Can be recognized by their suffix (ase) Hydrolase - Enzymes that facilitate the cleavage of bonds in molecules with the addition of the elements of water Activation Energy - Initial input of energy to start reaction - Allows molecules to get close enough to cause bond rearrangement - Can now achieve transition state where bonds are stretched - common ways to overcome activation energy - Large amounts of heat - Use of enzymes - How enzymes lower activation energy - Straining bonds in reactants to make it easier to achieve transition state - Position reactants together to facilitate bonding - Enzymes moving about until they collide at random Enzyme terminology - Active site: location where reaction takes place - Substrates: reactants that bind to active site - Enzyme substrate complex: formed when enzymes and substrate bind Substrate binding - Enzymes have a high specificity for their substrates - Lock and key metaphor for substrate and enzyme binding - only the right key will fit the lock (substrate and enzyme) - Induced fit phenomenon: interaction also involved conformational changes Enzyme reaction - Affinity: Degree of attraction between an enzyme and its substrate - Saturation: plateau where nearly all active sites are occupied by substrate  - Michaelis constant - Substrate concentration where velocity is half maximal value OR half of the active sites are occupied at one time - Substrate concentration required for the chemical reaction to occur - High enzymes needs higher substrate concentration - Inversely related to to affinity between enzyme and substrate Enzyme inhibitors - Competitive inhibition - Inhibitor molecules binds to active site - Inhibits ability of substrate to bind - needs to increase as more substrate is needed - Noncompetitive inhibition: lowers Vmax without affecting - Inhibitor binds to allosteric site, not the active site - Causes a shape change in the enzymes Other requirements for enzymes - Prosthetic group - Small molecules permanently attached to the enzyme and aids in enzymes function - Cofactor - Usually inorganic ion that temporarily binds to enzyme to promote a chemical reaction - Coenzyme: organic molecule that participates in reaction but is left unchanged afterward - Most enzyme function maximally in a narrow range of temperature and pH Overview of metabolism - Chemical reactions occur in metabolic pathways - Each step is coordinated by specific enzyme - Anabolic pathway - Synthesis cellular components - endergonic - Catabolic Pathways - Breakdown cellular components - Exergonic - Building large molecules (Anabolic)  - Building bigger molecules from smallers - Building cells and bodies - Proteins are synthesized by bonding amino acids Catabolics - Breakdown of reactants - Used for recycling building blocks - Used for energy to drive endergonic reactions - Energy stored intermediates such as NADH and ATP TWO ways to make ATP - Substrate level phosphorylation - Enzyme directly transfers phosphate from one molecules to another - Chemiosmosis - Energy stored in an electrochemical gradient is used to make ATP from ADP and P - Cellular respiration is metabolism Cellular respiration and ETC - Electron carriers, also called electron shuttles, are small organic molecules that play key roles in cellular respiration - Purpose is the movement of electrons across molecules - The reactions in which NAD+ and FAD gain or lose electrons are examples of a class of reactions called redox reactions - Name originates from oxidation-reduction reactions - Cellular respiration involves many reactions in which electrons get passed from one molecule to another - Oxidation is the removal of electrons - Reduction is the addition of electrons Regulation of metabolic Pathways - Needed for the cell to regulate materials - Catabolic pathways are regulated so that organic molecules are broken down ONLY AFTER they’re no longer needed or when the cell needs energy - Anabolic Pathways- ensures a cell synthesizes molecules when needed - Gene regulation: turn genes on or off that encode for the creation of enzymes - Cellular regulation: cell signaling pathways like hormones - Biomechanical regulation: feedback inhibition, product of pathway inhibits early steps to prevent over accumulation of product - Recycling of organic molecules - Most large molecules exist for a relatively short period of time - Half-life, time it takes for 50% of the molecules to be broken down and recycled - All living organisms must efficiently use and recycle organic molecules  - Expression of genomes Proteasome - Large complex that breaks down proteins using protease - Protease cleave bonds between amino acids - Ubiquitin tags target proteins - Degrades improperly folded proteins - Rapidly degrade proteins to respond to changing cell conditions Lysosomes - Lysosomes contain hydrolases to break down proteins, carbs NA’s, and lipids - Digest substances taken up by endocytosis - Autophagy: recycling worn out organelles using an autophagosome